Chemistry in the Public Domain (the author replies) - ACS Publications

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Chemical Education Today

Letters Chemistry in the Public Domain

Tetrahedral Bond Angle

I enjoyed the article by Sidney Toby in the November 1997 issue of this Journal (J. Chem. Educ. 1997, 74, 1285– 1287). I must take exception to the author’s assertion that ammonia–air mixtures are nonexplosive. During the period between my two teaching careers, I spent 17 years as a fire and explosion investigator, and several times found that simple answers are not always right. In this letter I will refer only to the literature that is on most chemistry teacher’s desks, not the fire and explosion profession’s literature. The Merck Index, 11th edition in article 510 states: “Mixtures of ammonia in air will explode when ignited under favorable conditions but ammonia is generally considered to be nonflammable.” Lange’s handbook, 11th edition in Table 11-10 lists the lower explosive limit for ammonia in air at 4.5%. There are many other references that support the fact that ammonia–air mixtures are explosive, but difficult to ignite. These two are available and familiar to most chemists. Chlorinated solvents are also a class of compounds that, while generally considered nonflammable, will also explode under favorable conditions. I once investigated a case where a chemist had assured an appliance manufacturer that methylene chloride was nonflammable and therefore its vapors were nonexplosive. A large (multistory) vapor degreaser exploded as a result of this ill-considered advice. I have experimentally exploded mixtures of both ammonia in air and methylene chloride in air.

I have been reading the Journal of Chemical Education and even contributed one paper four years ago (Ferreira, R. J. Chem. Educ. 1993, 70, 483). So I was pleased to learn that there are still people interested in how to calculate the tetrahedral bond angle (Glaister, P. J. Chem. Educ. 1997, 74, 1086). Fifty years ago last May, I published my first note in your magazine on this very problem (Ferreira, R. J. Chem. Educ. 1947, 24, 246). The most meaningful way to find the value of θ was shown to me by, of all people, Dick Feynman. Methane has no permanent electric dipole moment, but each C–H bond has one. Hence, any bond moment must be equal (but with opposite sign) to three other bond moments which make an angle θ, the cosine of which must be –1/3. Hence, θ = 109o28′.

John L. Odom Chattanooga School for the Arts and Sciences Chattanooga, TN 37403

The author replies: I thank John Odom for his comments. Before I wrote the paper, I put a lecture bottle of ammonia in a fume hood, opened the valve slightly and, somewhat timorously, put a lighted match in the ammonia stream several times. The flame flickered but whenever the match was removed, the flame went out. I was unable to make ammonia burn in the absence of the flame and I therefore support the U. S. Department of Transportation’s classification of it as a “nonflammable gas” under ordinary conditions. Can an ammonia–air mixture explode? Handbooks vary in their assessment but Bretherick’s Handbook (4th Ed., Butterworth) quotes 15.8% in air as a lower explosion or flammability limit. This is an order of magnitude higher than the value quoted in the article; nevertheless my statement that ammonia does not form potentially explosive fumes in air under ordinary conditions is probably wrong. Sidney Toby Department of Chemistry Rutgers University Piscataway, NJ 08855-0939

Ricardo Ferreira Department of Physics University of Calfornia, San Diego La Jolla, CA 92037-0345

Formation and Dimerization of NO2 I have strong misgivings about the article, “Formation and Dimerization of NO2” (J. Chem. Educ. 1997, 74, 1340– 1342). Aside from the unfortunate implication in the calculations that individual gases in a mixture occupy separate volumes, there is an overriding disadvantage to the experiment described: The results are terrible! A student who has any intuitive feel for quantitative relationships will be appalled by the scatter in the values obtained for the equilibrium “constant” and its distance from the accepted value. What appalls me, as one who spent 30 years trying to impart to students the faintest glimmer of the notion of “significant figures”, is the statement by the authors that 2.7±1.0 “compares reasonably well” with 8.6! More important, the beginning scientist may not know whether this is the result of poor data, or an innate characteristic of the scientific method. The implication that these are the kinds of data upon which the law of combining volumes is based will mislead the typical student, and confuse the circumspect one. The authors begin by observing that “experiments demonstrating…law of combining volumes are virtually absent” from general chemistry lab manuals. What their article illustrates most clearly is the reason for this observation: Successful performance of such an experiment is technically very demanding, and beyond the abilities of the typical general chemistry class. What would be the effect on a neophyte piano student of expecting him or her to play “The Minute Waltz” before ever having practiced scales? Discouragement and disillusion, surely, and perhaps the urge to try a different instrument! Edwin F. Meyer Emeritus Professor of Chemistry, DePaul University 1022 Dobson St. Evanston, IL 60202

JChemEd.chem.wisc.edu • Vol. 75 No. 9 September 1998 • Journal of Chemical Education

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